 In 1923, Bronsted and Lowry independently proposed a theory to explain the acid-base behavior of substances, in which acids are proton donors while bases are proton acceptors.

This theory naturally explains, unlike Arrhenius, the basic behavior of ammonia.

$NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-$

Water acts as an acid by donating a proton to ammonia, which in turn acts as a base. Water becomes its conjugate base; the hydroxide ion. While ammonia is converted to its conjugate acid; the ammonium ion.

This proton transfer between water and ammonia releases hydroxide ions into the medium, responsible for the basicity of the solution.

Since ammonia is a weak base, the reaction is not completely shifted towards the products, equilibrium being represented by a double arrow.

Equilibria of weak acids and weak bases are described by an equilibrium constant, called the acidity or basicity constant.

$K_b=\frac{[NH_4^+][OH^-]}{[NH_3]}$

Water acts as a solvent, its concentration not changing during the base dissociation process and is not included.

in the equilibrium constant.

If we consider the dissociation of acetic acid in water, we have:

$AcOH + H_2O \rightleftharpoons AcO^- + H_3O^+$

The ionization constant of the acid is given by:

$K_a=\frac{[AcO^-][H_3O^+]}{[AcOH]}$

In the case of strong acids, such as HCl, the dissociation is so important that they present a very high equilibrium constant (of the order of $10^6$), which makes it possible to treat the dissociation of this type of acid as complete. We indicate this fact by writing the ionization equation with a single arrow.

$HCl \rightarrow H^+ +Cl^-$